Average mass of an atom in unified atomic mass units (1 u is 1/12 of the mass of a carbon-12 atom). Accounts for the natural mix of the element’s isotopes.
Number of the vertical column, 1 to 18 in IUPAC numbering. Elements in a group share the same outer-shell layout, so their properties are alike.
Number of the horizontal row. Usually equal to the number of electron shells in the atom; palladium is the exception, with four shells in period 5.
The subshell (s, p, d or f) that receives the last electrons in the usual filling order; it divides the table into four blocks. Real atoms sometimes differ: lanthanum starts the f-block but has no f electrons.
[Ne] 3s² 3p²
How electrons are arranged in subshells. The bracket stands for the previous noble gas; the superscript is the number of electrons in a subshell.
2, 8, 4
Number of electrons in each shell, from the nucleus outwards.
1,414 °C (1,687 K)
Temperature at which the solid turns liquid at standard pressure. Helium, carbon and arsenic do not melt at normal pressure: helium needs about 25 atm, carbon and arsenic turn straight into vapour.
3,265 °C (3,538 K)
Temperature at which the liquid boils at standard pressure. For carbon and arsenic, the temperature at which the solid turns straight into vapour (sublimation).
2.33 g/cm³
Mass of one cubic centimetre of the substance. For gases, at 0 °C and standard pressure.
1.90 Pauling
How strongly an atom pulls electrons towards itself in a chemical bond. On the Pauling scale it runs from 0.79 for caesium to 3.98 for fluorine.
8.152 eV
Energy needed to remove the most loosely bound electron from an atom. The lower it is, the more easily the element gives up an electron, which is typical of metals.
1.39 eV
Energy released when a neutral atom gains an electron. Highest for the halogens. Noble gases, nitrogen and some metals form no stable negative ion.
111 pm
Half the length of a single covalent bond between two atoms of the element. It shows the size of an atom inside a molecule or crystal: 76 pm for carbon, 132 pm for iron.
+4, +2, -4
Charges the atom would have in its compounds if all its bonds were ionic; the main ones are listed.
Silicon (Si) is a chemical element with atomic number 14. It sits in group 14 and period 3 of the periodic table, in the p-block. It is a metalloid. At room temperature it is a solid. Its standard atomic weight is 28.085.
Origin of the name
The name comes from Latin silex, ‘flint’, a stone made of silica. Humphry Davy proposed silicium in 1808, with the -ium ending of metals, because he took it for one. In 1817 Thomas Thomson changed the English name to silicon, to match carbon and boron, which silicon resembles.
Where is silicon found?
Silicon is the second most abundant element in the Earth’s crust after oxygen, more than a quarter of its mass. It never occurs free, always bound to oxygen: as silica in sand, quartz, flint, amethyst, agate and opal, and as silicates in feldspar, mica, clay and granite.
What is silicon used for?
Most silicon is used without being purified: sand, clay and stone become concrete, bricks and glass. Silicon is added to steel and to aluminium alloys. Purified to a very high degree, it is the material of computer chips, transistors and solar cells. It is also the backbone of silicones, used in sealants, lubricants and medical implants, and silicon carbide is a hard abrasive.
Interesting facts
Silicon Valley in California is named after the silicon chips made there.
Tiny algae called diatoms build their shells out of silica.
People in Egypt were making glass from silica sand at least 3,500 years ago.